Calculate the value of ()Go for the following reaction , given the standard free energies of formation listed below each ion.
Ag(SO3)_2^3- -> Ag+ + 2SO32-
delta Gf, kJ/mol = -943 +77 -497
a. - 1860 kJ b. – 128 kJ c. + 26 kJ d. - 369 kJ e. + 523 kJ
help no idea
dGrxn = (n*dGproducts) - (n*dGreactants)
I don't think so. Show your work.
not getting right answer how do you set it up?
I'm doing all the work. Show your work and I will know what you're doing wrong. I think you are guessing.
is this correct?
(1 x -943) - (2 x -497) + 77
-943 - 917 = -1860 so A?
Calculate the value of ()Go for the following reaction , given the standard free energies of formation listed below each ion.
Ag(SO3)_2^3- -> Ag+ + 2SO32-
delta Gf, kJ/mol = -943 +77 -497
a. - 1860 kJ b. – 128 kJ c. + 26 kJ d. - 369 kJ e. + 523 kJ
is this correct?
(1 x -943) - (2 x -497) + 77
-943 - 917 = -1860 so A?
To calculate the standard Gibbs free energy change (ΔG°) for a reaction, you can use the equation:
ΔG° = ∑(ΔGf° products) - ∑(ΔGf° reactants)
where ΔGf° is the standard free energy of formation.
Given the standard free energies of formation for each ion involved in the reaction, we can substitute the values into the equation.
Reactants:
Ag(SO3)2^3- ΔGf° = -943 kJ/mol
Products:
Ag+ ΔGf° = +77 kJ/mol
2SO32- ΔGf° = -497 kJ/mol
Now, let's substitute these values into the equation:
ΔG° = [(+77 kJ/mol) + 2(-497 kJ/mol)] - (-943 kJ/mol)
Simplifying:
ΔG° = 77 kJ/mol - 994 kJ/mol
ΔG° = -917 kJ/mol
So, the value of ΔG° for the reaction is -917 kJ/mol.
None of the options provided match the calculated value of -917 kJ/mol, so it seems there might be an error in the given options.